Sulfur trioxide, a precursor of sulfuric acid, is prepared industrially using the contact process by passing sulfur dioxide and oxygen over a hot catalyst:

Question

2SO2(g) + O2 (g) <--> 2SO3 (g)
Using a table of thermodynamic data, estimate the temperature at which the reaction begins to favor the formation of products at equilibrium. Round your answer to 3 significant digits.

Answer 1

To determine the temperature at which the reaction begins to favor the formation of products at equilibrium, we need to calculate the equilibrium constant (K) for the reaction using the equation:

ΔG° = -RT ln(K)

First, we need to calculate the standard Gibbs free energy change (ΔG°) for the reaction using the following equation:

ΔG° = ΣnG°(products) - ΣnG°(reactants)

Using standard Gibbs free energy of formation values from tables:

ΔG° = 2(-396.2 kJ/mol) - 0 kJ/mol - (2(-296.8 kJ/mol)) = -198.8 kJ/mol

Now, we convert the ΔG° to joules:

ΔG° = -198.8 kJ/mol * 1000 J/kJ = -198,800 J/mol

Now, we can rearrange the equation to solve for the equilibrium constant (K):

K = e^(-ΔG°/RT)

We don't have the exact value for R (gas constant), but it's usually taken as 8.314 J/(mol*K). We can assume a temperature of 298 K to get an estimate.

K = e^(-(-198,800 J/mol)/(8.314 J/(mol*K) * 298 K))

K ≈ e^(252.93) ≈ 3.4 x 10^109

So, at a temperature of around 298 K, the reaction begins to favor the formation of products at equilibrium.